1. Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases
Chem 106, Chapter 17
Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases
Copyright 2007, David R. Anderson

2. Acids and Bases
Arrhenius definition
Brønsted definition
Relative strengths of conjugate acid-base pairs
Effect of structure on acid-base strength
Acid-base equilibria
Definitions
pH, pOH
Ka, Kb, pKa, pKb
Kw: autodissociation of water
pH Calculations
Strong acids and bases
Weak acids and bases
Polyprotic acids
Salts
Lewis Acids and Bases

3. I. Acids and Bases A. Arrhenius definition B. Brønsted definition
produces hydronium ion (H3O+) in aqueous solution
produces hydroxide ion (OH–) in aqueous solution
B. Brønsted definition
acid:
base:
donates a proton (hydrogen ion, H+)
accepts a proton
HF H2O
H3O F–
A conjugate acid is formed by adding a proton to something.
A conjugate base is formed by removing a proton from something.
acid
base
conjugate
acid
conjugate
base
NH H2O
NH OH–
base
acid
conjugate
acid
conjugate
base

4. I. Acids and Bases B. Brønsted definition F– + H2O HF + OH– NH4+ + H2O
water is amphoteric or amphiprotic (behaves as either an acid or a base)
H2O H2O
H3O OH–
base
acid
autoionization of water:
Brønsted definition fits even when not in aqueous solution:
NH HF
NH F–
base
acid

5. I. Acids and Bases C. Relative strengths of conjugate acid-base pairsHA
H3O+ + A–
 If HA is a stronger acid then A– is a weaker base.
If HA is a weaker acid then A– is a stronger base. B
BH+ + OH–
 If B is a stronger base then BH+ is a weaker acid.
If B is a weaker base then BH+ is a stronger acid.

6. I. Acids and Bases D. Effect of structure on acid-base strength
General: HA
H3O+ + A–
all acids
produce this
the place to look for
a difference is here
The strength of HA depends on the stability of A–.
If A– is more stable
(weaker base)
 HA will be a stronger acid
If A– is less stable
(stronger base)
 HA will be a weaker acid
A– is more stable when: • A is more electronegative
• the negative charge is more delocalized

7. I. Acids and Bases D. Effect of structure on acid-base strength
1. binary acids: hydrogen + one other element
acidity
increases
(a)
(b) HA
H3O+ + A–
(a) As A becomes more electronegative, A– is more stable (weaker base),
and HA is a stronger acid.
Thus: HF > H2O > NH3 > CH4
(dissociates: 1% –7% –16% –23%)
Similarly: HCl > H2S > PH3 > SiH4

8. I. Acids and Bases D. Effect of structure on acid-base strength
1. binary acids: hydrogen + one other element
acidity
increases
(a)
(b) HA
H3O+ + A–
(b) As A– becomes larger, the negative charge is more delocalized,
A– is more stable (weaker base), and HA is a stronger acid.
(overcomes opposite trend in electronegativity)
Thus: HI > HBr > HCl > HF
strong acids
(100% dissociated)
weak acid
(1% dissociated)
Similarly: H2Te > H2Se > H2S > H2O

9. I. Acids and Bases D. Effect of structure on acid-base strength
2. oxoacids: HnXOm (HNO3, H2SO4, H3PO4, etc.)
a. electronegativity of central atom (with same number of oxygens)
e.g., HClO4 is a stronger acid than HBrO4
central atom withdraws electron density from the oxygen, stabilizes the anion
Cl is more electronegative than Br
 more stable anion, stronger acid
Thus: HClO4 > HBrO4 > HIO4
HClO3 > HBrO3 > HIO3
HClO2 > HBrO2 > HIO2
HClO > HBrO > HIO
And: HClO4 > H2SO4 > H3PO4
HBrO4 > H2SeO4 > H3AsO4
etc.

10. I. Acids and Bases D. Effect of structure on acid-base strength
2. oxoacids: HnXOm (HNO3, H2SO4, H3PO4, etc.)
b. number of lone oxygens (with same central atom)
negative charge more delocalized,
anion more stable  stronger acid
Thus: HClO4 > HClO3 > HClO2 > HClO
HBrO4 > HBrO3 > HBrO2 > HBrO
etc.
And: H2SO4 > H2SO3 H2SeO4 > H2SeO3
HNO3 > HNO2 H3PO4 > H3PO3

11. I. Acids and Bases D. Effect of structure on acid-base strength
Chem 106, Chapter 17
I. Acids and Bases
D. Effect of structure on acid-base strength
3. carboxylic acids
(10-6 % dissociated)
(~1 %
dissociated)
negative charge delocalized,
anion more stable, stronger acid
Copyright 2007, David R. Anderson

12. I. Acids and Bases D. Effect of structure on acid-base strength
Which in each pair would be the stronger acid?
H2SO3 or HClO3 CH3OH or CH3SH
HOBr or HBrO3 HBrO2 or HClO3
F2CHCO2H or FCH2CO2H ClCH2CH2CO2H or CH3CHClCO2H
HNO3 or HNO2 PH3 or NH3
HClO4 or HBrO4 H2O or HF
H2S or H2Se FCH2CO2H or ClCH2CO2H
H2Se or HBr HF or HCl

13. I. Acids and Bases Summary: acid and base strength
“Absolute” strengths (in water): HA
H3O+ + A–
strong acid nonbasic K ~  (100% dissociated)
weak acid weak base K ~ 10–3 - 10–10
neutral strong base K ~ 10–15
nonacid v. strong base K 10–20 ~
<

14. I. Acids and Bases
Summary: acid and base strength

15. I. Acids and Bases E. Acid-base equilibria HA + B BH+ + A–
Chem 106, Chapter 17
I. Acids and Bases
E. Acid-base equilibria
HA B
BH A–
If: stronger stronger weaker weaker then K > 1
acid base acid base (equilibrium lies on right)
On which side does the equilibrium lie in the following reactions?
OCl– + HCl
NH ClO4–
F– + H2O
NH2– + H2O
Copyright 2007, David R. Anderson

16. II. Definitions A. pH, pOH Definition: pX = -logX
pH = -log[H3O+] pOH = -log[OH–]
[H3O+] = 10–pH [OH–] = 10–pOH
[H3O+] pH
0.01 M
1 x 10–7 M
3 x 10–10 M
pH [H3O+]
3.0
8.4

17. II. Definitions A. pH, pOH sig figs in pH calculations:
e.g., [H3O+] = 2.3 x 10–3 M
log [H3O+] = log
2 sig figs
exact
= ….
log [H3O+] = -2.64
pH = -log [H3O+] = 2.64
from exponent
# of sig figs in [H3O+]
 # of sig figs in [ ] = # of places past decimal in pH

18. II. Definitions B. Ka, Kb, pKa, pKb HA + H2O H3O+ + A– [H3O+][A–]
Kc =
~ constant (55 M)
 Kc·[H2O] = Ka =
[H3O+][A–]
[HA]
acid dissociation constant
pKa = -log Ka if pKa = 5 then Ka = 10–5
if pKa = 8 then Ka = 10–8
stronger acid
weaker acid

19. II. Definitions B. Ka, Kb, pKa, pKb B + H2O BH+ +OH– [BH+][OH–]
Kc =
~ constant
 Kc·[H2O] = Kb =
[BH+][OH–]
[B]
base dissociation constant
pKb = -log Kb if pKb = 4 then Kb = 10–4
if pKb = 9 then Kb = 10–9
stronger base
weaker base

20. II. Definitions C. Kw: autodissociation of water H2O + H2O H3O+ + OH–
Kw = [H3O+][OH–] = 10–14 (constant at 25ºC)
 [H3O+] =
10–14
[OH–]
pKw = pH + pOH = 14
pH = 14 - pOH

21. III. pH Calculations A. Strong acids and bases 100% dissociated
 for strong acid: [H3O+]eq = [HA]I
base: [OH–]eq = [B]i
e.g., 1.0 x 10–3 M HCl
[H3O+] = pH =
[OH–] = pOH =
e.g., 2.5 x 10–2 M NaOH
[OH–] = pOH =
[H3O+] = pH =
[OH–]
-1.60
+ 14
log
pH = 12.40

22. III. pH Calculations B. Weak acids and bases [H3O+][A–] [HA] HA
Solve equilibrium
expressions
[BH+][OH–]
[B] B
BH+ + OH–
Kb =
e.g., What is the pH of 0.10 M HC2H3O2? (Ka = 1.8 x 10–5)
HC2H3O2
H3O+ + C2H3O2–
1.8 x 10–5 = x2
( x)
x = [H3O+] = 1.3 x 10–3 M
(assumption valid)
pH = 2.87
assume x << 0.10
1.8 x 10–5 = x2
(0.10)
 If [HA]i  400·Ka then x << [HA]i
(If not, then have to solve quadratic.)

23. III. pH Calculations B. Weak acids and bases
e.g., What is the pH of 0.25 M N2H4? (Kb = 1.7 x 10–6)
N2H4

24. III. pH Calculations B. Weak acids and bases
e.g., If a 0.10 M solution of a weak acid has a pH of 4.28, what is the Ka for the acid?

25. III. pH Calculations B. Weak acids and bases
e.g., What concentration of C5H5N (Kb = 1.7 x 10–9) will give a solution with pH = 9.80?
C5H5N

26. III. pH Calculations C. Polyprotic acids H2SO4, H2SO3, H2CO3, etc.
Lose their protons in separate steps:
H2A
H3O+ + HA–
Ka1 =
[H3O+][HA–]
[H2A]
(Usually, Ka1 >> Ka2)
HA–
H3O+ + A2–
Ka2 =
[H3O+][A2–]
[HA–]
Assume: 1) [H2A], [H3O+], and [HA–] can be determined from the 1st step.
(i.e., HA– dissociates only very little.)
2) [A2–] can be determined from the 2nd step.

27. III. pH Calculations C. Polyprotic acids
e.g., What are the equilibrium concentrations of all species in 0.10 M Vitamin C (ascorbic acid, H2C6H6O6)?
H2A
H3O+ + HA–
Ka1 = 7.9 x 10–5
HA–
H3O+ + A2–
Ka2 = 1.6 x 10–12
Ka1 =
[H3O+][HA–]
[H2A]
Ka2 =
[H3O+][A2–]
[HA–]
7.9 x 10–5 = x2
( x)
(0.10) = ~
1.6 x 10–12 =
(2.8 x 10–3 + y)(y)
(2.8 x 10–3 - y) = ~
(2.8 x 10–3)(y)
(2.8 x 10–3)
= y
x = 2.8 x 10–3 M
y = 1.6 x 10–12 M
 [H2A] = 0.10 M
[H3O+] = [HA–] = 2.8 x 10–3 M
[A2–] = 1.6 x 10–12 M

28. III. pH Calculations C. Polyprotic acids
e.g., What are the equilibrium concentrations of all species in a 0.10 M solution of carbonic acid, H2CO3? (Ka1 = 4.2 x 10-7, Ka2 = 4.8 x 10-11)

29. III. pH Calculations D. Salts HA + B  BH+A– (salt)
Salt may be acidic or basic depending on the nature of BH+ and A–.
conjugate acid: BH+ + H2O
B + H3O+
conjugate base: A– + H2O
HA + OH–

30. III. pH Calculations D. Salts Ka and Kb for conjugate acid-base pairs:
[H3O+][A–]
[HA] HA
H3O+ + A–
Ka = A–
HA + OH–
Kb =
[HA][OH–]
[A–]
usually only one or the other given in a table
Ka · Kb =
[H3O+][A–]
[HA]
[HA][OH–]
[A–]
= [H3O+][OH–] = Kw
or Ka =
or Kb = Kw Ka Kb
Ka · Kb = Kw
for a conjugate
acid-base pair

31. III. pH Calculations D. Salts 1. acidic salts
e.g., What is the pH of a 0.10 M solution of NH4Cl?
NH4+:
Cl–:
conjugate acid of a weak base  weak acid
conjugate base of a strong acid  nonbasic (spectator)
NH4+
NH3 + H3O+
Ka(NH4+) = Kw
Kb(NH3) =
[H3O+][NH3]
[NH4+]
10–14
1.8 x 10–5 x2
( x) x2
(0.10) = = ~
x = [H3O+] = 7.5 x 10–6 M pH = 5.13

32. III. pH Calculations D. Salts 2. basic salts
e.g., What is the pH of a 0.10 M solution of NaCN? Ka(HCN) = 4.9 x 10–10
Na+:
CN–:

33. IV. Lewis Acids and Bases
acid: accepts a pair of electrons
base: donates a pair of electrons
HBr + H2O
H3O+ + Br–
Arrhenius,
Brønsted,
Lewis
NH3 + H2O
NH OH–
Arrhenius,
Brønsted,
Lewis

34. IV. Lewis Acids and Bases
HBr + NH3
NH4+ Br–
Brønsted,
Lewis
NaI + CH3Br
CH3I + NaBr
Lewis
only

35. IV. Lewis Acids and Bases
NaOH + CO2
NaHCO3
Lewis
only